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Gain of electrons, removal of oxygen atoms or addition of hydrogen atoms

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Loss of electrons, addition of oxygens or removal of hydrogen

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Advantages:

A convenient way of storing & releasing energy

energy efficiency is much higher than standard fuel systems(36-45% compared to diesel 22%)

emissions are less damaging than CO₂ from traditional engines

Disadvantages:

gases needed are difficult to store compared to liquid fuel

fuel cells operate at lower temps(80°C) so need efficient catalysts which sue expensive metals

hydrogen fuel must be generated elsewhere & is likely to use fossil fuels causing CO₂ emissions and the process is not 100% efficient

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Currently being developed as a method of efficient energy release from fuels such as hydrogen, methane or methanol.  An electrochemical method of releasing energy. The fuel cell system passes the fuel over platinum metal which acts as a catalyst, but also as an electrode for the electrochemical system. Electrons are removed from hydrogen at 1 electrode
H₂ + 2H⁺ --> 2e-
The protons(H+) diffuse through a semi-permeable membrane to the other electrode where they receive oxygen molecules, creating water.
O₂ + 4H+ + 4e⁻ --> 2H₂O
The overall reaction is 2H₂ + O₂ --> 2H₂O and the overall voltage is 1.23V

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A feasible reaction is one that occurs spontaneously. The oxidising agent must be strong enough to oxides the reducing agent.
For the reaction to go ahead:

the EMF value must be positive

tho oxidation half-reaction must be more negative

the reduction reaction must be more positive

if you have 2 negative half-reactions, the reduction reaction is the least negative one

when calculating EMF you must work out which half-reaction is the more positive(reduction) and more negative(oxidation)

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The EMF value is given by the difference between the standard electrode potential of the two half-cells.
E θ cell = E θ (most positive/reduction) - E θ (most negative/oxidation)

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The positive electrode has the more positive E θ

The electrons flow to the half-cell with the more positive E θ

Using this method we can put the reducing power of any set of half-cells in order(electrochemical series). This order is always shown as reduction reactions, so you always see ions as gaining electrons. It's noticeable that the order for metals reflects the reactivity series. The most reactive metals have the most negative E θ value and the most reactive non-metal has the most positive E θ value.
The most oxidising agent has the highest positive value and the most reducing agent has the most negative value.

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Consisting of a platinum electrode coated with fine platinum grains, called platinum black. H+ ions as hydrochloric acid have hydrogen gas passed over it slowly. To find the standard electrode potential of a half-cell it is joined to the hydrogen electrode creating a system.
The standard conditions are:

1 ATM of pressure

298K(25°) temperature

1 mol dm⁻³ concentration

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Using half-cells allows us to gain detailed info about the reaction, by calculating the EMF of the reaction we can measure the difference in redox power of the 2 halves. The largest EMF value will be of the easily oxidised species & the lower value for the easily reduced species.
The ability for a half-cell to gain/lose electrons is measure using Eθ, the scale uses hydrogen as 0 & any species that is easier to reduce has a negative Eθ value, with easier oxidising species having a more positive value. The hydrogen half-cell is called the standard hydrogen electrode.

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If a question asks for a cell diagram, this is what is expected.
Starting with the conducting metal, if there is a change of state then a single vertical line is in-between the 2 ions. If the 2 ions are the same state then a comma is used. Then a double vertical line represents the salt bridge. Then the same for the other half-cell. The most positive electrode is placed on the right-hand side.

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Solution of a metal in 2 different oxidation states

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Solution of a metal in 2 different oxidation states

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There is no conductor so we must use an inert metal electrode. This is typically used for transition metals, where the metal may have several oxidation states. Both Fe²⁺/Fe3+ and Mn²⁺/MnO4-. Both cause colour change when reduction or oxidation occurs. Fe²⁺ is pale grey, Fe³⁺ is yellow/orange; Mn²⁺ is colourless and MnO4- is purple.

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Non-metals aren't conductors so an inert platinum electrode allows electrons to flow in/out of the half-cell. This is typically used for hydrogen. The gas is bubbled over the inert electrode which is in a solution of ions. The changes do not cause any apart colour change.

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Two examples so far are Zn(s) & Zn²⁺(aq) and the Cu(s) & Cu²⁺(aq) half-cells. In both we have a metal electrode with a solution containing a 1moldm⁻³ solution of the metal ions. In the case of zinc, there is no apparent colour change,  but for copper the blue solution may lose colour as the copper ions are reduced.

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Each cell is made up of 2 half-cells, each of which must contain both the reactants & products of the half-equation. Meaning the half-cell for the reduction of Cu2+ ions to copper atoms must contain both of these. It must contain a metal to allow electrons to flow into or out of the half-cell. There are various types of half-cell depending on the physical states of each substance in reaction.

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Practically it is possible to separate 2 processes so oxidation happens in one place & reduction in another. To do this we set up two half-cells, joining them together to create a complete circuit.

Wire allows electrons to flow from HC where oxidation occurs to HC where reduction occurs. Often including a high resolution voltmeter if we’re measuring the potential difference in the cell

Salt bridge completes the circuit & allows ions to flow without the solutions mixing. A typical salt bridge is made of a gel soaked in a solution of potassium nitrate. A simpler salt bridge is made of filter paper soaked in the same solution

The entire apparatus is called a cell, with the 2 parts called a half-cell

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Electrons are transferred from one species to another; a combination of reduction & oxidation.
OIL RIG!!

An oxidising agent is a species that oxidises another, & itself becomes reduced.

A reducing agent is a species that reduces another, & itself becomes oxidised.

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Uncombined element - 0

sum of oxidation numbers with no charge - 0

fluorine - -1

oxygen - -2 except in peroxides(-1)

hydrogen - +1 except in metal hydrides(-1)

group 1 metals - +1

group 2 metals - +2

halogen in metal halide - -1